1. Dr. Vatsala Soni PGGC Sector -11 Chandigarh
The s-Block Elements
Dr. Vatsala Soni
PGGC Sector Chandigarh

2. Members of the s-Block ElementsLi Be Na K Rb Cs Fr Mg Ca Sr Ra Ba
IA IIA
IA Alkali metals
IIA Alkaline Earth
metals

3. Chapter summary Characteristic properties of the s-block elements
Variation in properties of the s-block elements
Variation in properties of the s-block compounds
Uses of compounds of the s-block elements

4. Characteristic properties of s-block elements
Metallic character
Low electronegativity
Basic oxides, hydroxides
Ionic bond with fixed oxidation states
Characteristic flame colours
Weak tendency to from complex

5. Metallic character High tendency to lose e- to form positive ions
Metallic character increases down both groups

6. Electronegativity Low nuclear attraction for outer electrons
Highly electropositive
Small electronegativity
Group I
Group II Li Be Na
Mg 1.2 K Ca Rb Sr Cs Ba Fr Ra

7. Basic oxides, hydroxides
Li2O
LiOH
Na2O, Na2O2
NaOH
K2O2, KO2
KOH
Rb2O2, RbO2
RbOH
Cs2O2, CsO2
CsOH
Oxide
Hydroxides
BeO
Be(OH)2
MgO
Mg(OH)2
CaO
Ca(OH)2
SrO
Sr(OH)2
BaO, Ba2O2
Ba(OH)2

8. Oxides, Peroxide, Superoxide
Reaction with water:
Oxide: O2- + H2O  2OH-
Peroxide: O H2O  H2O2 + 2OH-
Superoxide: 2O2- + 2H2O  2OH- + H2O2 + O2
:O:.O:
.. ..
Super oxide
:O:O:
.. ..
Peroxide ion
Li does not form
peroxide or super oxide
Li2O2  Li2O + ½ O2

9. Hydroxides Group I hydroxides Li Na K Rb Cs
All are soluble, base strength
increase.
Group II
hydroxide Be Mg Ca Sr Ba
Solubility increase, from
Amphoteric to basic, base strength
increase

10. Predominantly ionic with fixed oxidation state
Group I: Most electropositive metals.
Low first I.E. and extremely high second I.E.
Form predominantly ionic compounds with
non-metals by losing one electron.
Fixed oxidation state of +1.
Group II: Electropositive metals.
Low first and second I.E. but very high third
I.E Have a fixed oxidation state of +2.
Be and Mg compounds possess some degree
of covalent character.

11. Characteristic flame colours
Na+ Cl- (g)  Na (g) + Cl (g)
Na(g)  Na* (g)
[Ne]3s [Ne]3p1
Na*(g)  Na(g) + h (589nm, yellow)

12. Flame test Li deep red Ca brick red Na yellow Sr blood red K lilac
Rb bluish red
Cs blue
Ca brick red
Sr blood red
Ba apple green
HCl(aq)
sample

13. Weak tendency to form complex
Complex formation is a common feature of d-block
element. e.g. Co(NH3)63+ Co
:NH3
H3N:
s-block metal ions have no low energy vacant
orbital available for bonding with lone pairs
of surrounding ligands, they rarely form complexes.

14. Variation in properties of elements
Atomic radii
Ionization enthalpies
Hydration enthalpies
Melting points
Reactions with oxygen, water, hydrogen and chlorine

15. Atomic radii (nm) Li 0.152 Be 0.112 Na 0.186 Mg 0.160 K 0.231 Ca 0.197Rb
0.244 Sr
0.215 Cs
0.262 Ba
0.217 Fr
0.270 Ra
0.220 Fr Ra Li Be

16. Ionization Enthapy Group I 1st I.E. 2nd I.E. Li 519 7300 Na 494 4560 K
418
3070 Rb
402
2370 Cs
376
2420
Group I
1st I.E.
2nd I.E.
3rd I.E. Be
900
1760
14800 Mg
736
1450
7740 Ca
590
1150
4940 Sr
548
1060
4120 Ba
502
966
3390

17. Ionization Enthalpy 1st I.E. Be+ Li Na 2nd IE K Rb Ca+ Cs Ba+ Be Ca Ba
300
400
500
600
500
1000
1500
2000 Be Ca Ba
Be+
Ca+
Ba+
1st IE
2nd IE

18. Ionization Enthalpy Group I
Have generally low 1st I.E. as it is well shielded from the nucleus by inner shells.
2. Removal of a 2nd electron is much more difficult because it involves the removal of inner shell electron.
3. I.E. decreases as the group is descended.
As atomic radius increases, the outer e is further away from the well-shielded nucleus.

19. Ionization Enthalpy Group II Have low 1st and 2nd IE.
Removal of the 3rd electron is much more difficult
as it involves the loss of an inner shell electron.
IE decrease as the group is descended.
IE of the group II is generally higher than group I.

20. Hydration Enthalpy M+(g) + aqueous  M+(aq) + heat M+ -600 -300
Li+ Na+ K+ Rb+ Cs+ M+

21. Hydration Enthalpy Be2+ Mg2+ Ca2+ Sr2+ Ba2+ -2250 -2000 -1750 -1500
-600
-300
Li+ Na+ K+ Rb+ Cs+

22. Hydration Enthalpy General trends:
On going down both groups, hydration enthalpy
decreases.
(As the ions get larger, the charge density of the
ions decreases, the electrostatic attraction between
ions and water molecules gets smaller.)
Group 2 ions have hydration enthalpies higher
than group 1.
( Group 2 cations are doubly charged and have
smaller sizes)

23. Variation in Melting Points
250
500
750
1000
1250 Be Mg Ca Sr Ba Li Na K Rb Cs

24. Variation in Melting Points
Strength of metallic bond depends on:
Ionic radius
Number of e- contributed to the electron sea per atom
Crystal lattice structure
Note: The exceptionally high m.p. of calcium
is due to contribution of d-orbital participation
of metallic bonding.

25. Variation in Melting Points
Group I
Structure
Group II Li
B.C.C. Be
H.C.P. Na Mg K Ca
C.C.P. Rb Sr Cs Ba

26. Reactions with oxygen S-block elements are strong reducing agents.
Their reducing power increases down both groups.
(As the atomic size increases, it becomes easier to
remove the outermost electron)
S-block elements reacts readily with oxygen.
Except Be and Mg, they have to be stored under
liquid paraffin to prevent contact with the atmosphere.

27. Reactions with oxygen Normal Oxide Peroxide Superoxide Structure
Formed by
Li and Group II
Na and Ba
K, Rb, Cs
.. 2-
:O: ..
:O-O:
.. ..
:O:.O:
.. ..

28. Reaction with water M(s)  M+(aq) + e- H2O(l) + e-  OH-(aq) + ½ H2(g)
Li volt Na K Rb Cs
Be volt
Mg -2.38 Ca Sr Ba
Energetic vs. Kinetic Factor

29. Reaction with hydrogen
All the s-block elements except Be react directly with
hydrogen.
2Na(s) + H2(g)  2NaH(s)
Ca(s) + H2(g)  CaH2(s)
The reactivity increases down the group.
Only BeH2 and MgH2 are covalent, others are ionic.

30. Reaction with chlorine
All the s-block metals react directly with chlorine
to produce chloride.
All group I chlorides are ionic.
BeCl2 is essentially covalent, with comparatively low
m.p.
The lower members in group II form essentially ionic
chlorides, with Mg having intermediate properties.

31. Although lithium has highly negative Eo, it only
reacts slowly with water. This illustrates the importance
of the role of kinetic factors in determining the rate
of a chemical reaction.
Lithium has a higher m.p., this increases the activation
energy required for dissolution in aqueous solution.
It does not melt during the reaction as Na and K do, and
thus it has a smaller area of contact with water.

32. Variation in properties of the compounds
Reactions of oxides and hydroxides
Reactions of chlorides
Reactions of hydrides
Relative thermal stability of carbonates and hydroxides
Relative solubility of sulphate(VI) and hydroxde

33. Reactions of oxides and hydroxides
All group I oxides reacts with water to form
hydroxides
Oxide: O2- + H2O  2OH-
Peroxide: O H2O  H2O2 + 2OH-
Superoxide: 2O2- + 2H2O  2OH- + H2O2 + O2
All group I oxides/hydroxides are basic and the
basicity increases down the group.

34. Reactions of oxides and hydroxides
Group II oxides/hydroxides are generally less basic
than Group I. Beryllium oxide/hydroxide are
amphoteric.

35. Reactions of chlorides
All group I chlorides are ionic and readily
soluble in water. No hydrolysis occurs.
Group II chlorides show some degree of covalent
character.
Beryllium chloride is covalent and hydrolysis to
form Be(OH)2(s) and HCl(aq).
Magnesium chloride is intermediate, it dissolves and
hydrolysis slightly.
Other group II chlorides just dissolve without
hydrolysis.

36. Reactions of hydrides They all react readily with water to give the
metal hydroxide and hydrogen due to the
strong basic property of the hydride ion, H:-
H:-(s)+ H2O(l)  H2(g)+ OH-(aq)
Hydride ions are also good reducing agent.
They can be used to prepare complex hydrides
such as LiAlH4 and NaBH4 which are used to
reduce C=O in organic chemistry.

37. Thermal Stability Thermal stability refers to decomposition of the
compound on heating. Increased thermal stability
means a higher temperature is needed to decompose
the compound.

38. Thermal Stability of carbonates
Li2CO3  Li2O + CO2 ( at 700oC)
All other group I carbonates are stable at ~800oC
BeCO3  BeO + CO2 ( at 100oC)
MgCO3  MgO + CO2 ( at 540oC)
CaCO3  CaO + CO2 ( at 900oC)
SrCO3  SrO + CO ( at 1290oC)
BaCO3  BaO + CO2 ( at 1360oC)

39. Thermal Stability of hydroxides
All group I hydroxides are stable except LiOH
at Bunsen temperature.
Be(OH)2(s)  BeO(s) + H2O(g) H = +54 kJ/mol
Mg(OH)2(s)  MgO(s) + H2O(g) H = +81 kJ/mol
Ca(OH)2(s)  CaO(s) + H2O(g) H = +109 kJ/mol
Sr(OH)2(s)  SrO(s) + H2O(g) H = +127 kJ/mol
Ba(OH)2(s)  BaO(s) + H2O(g) H = +146 kJ/mol

40. Thermal stability Carbonates and hydroxides of Group I metals
are as a whole more stable than those of Group II.
Thermal stability increases on descending the group.
Lithium often follow the pattern of Group II rather
than Group I.
This is an example of the diagonal relationship.

41. Explanation of Thermal Stability
Charge of the ions
Size of the ions
Compounds are more stable if the charge increases
and size decreases.
For compounds with large polarizable anions, thermal
stability is affected by the polarizing power of the
cations.

42. Explanation of Thermal Stability- +
Decreasing
polarizing
power +
Increasing
stability - + -

43. Explanation of Thermal Stability
Mg2+ C O
O:-
- :O
Mg2+ O2- + CO2
Mg2+
Mg2+ O2- + H2O
-:O H

44. Explanation of Thermal Stability
MgCO3
MgO
MgO
BaO
BaCO3
BaO

45. Relative solubility of Group II hydroxides
Compound
Solubility / mol per 100g water
Mg(OH)2
0.020 x 10-3
Ca(OH)2
1.5 x 10-3
Sr(OH)2
3.4 x 10-3
Ba(OH)2
15 x 10-3
Solubility of hydroxides
increases down the group.

46. Solubility of Group II sulphates
Compound
Solubility / mol per 100g water
MgSO4
3600 x 10-4
CaSO4
11 x 10-4
SrSO4
0.62 x 10-4
BaSO4
0.009 x 10-4
Solubility of sulphates
increases up the group.

47. Explanation of solubility
MX(s)
aqueous
H solution
M+(aq) + X-(aq)
M+(g) + X-(g)
H hydration
-H lattice
H solution
-H lattice
H hydration = +

48. Explanation of solubility
Group I compounds are more soluble than Group II
because the metal ions have smaller charges and
larger sizes. H lattice is smaller, and H solution is
more exothermic.
H solution
-H lattice
H hydration = +

49. Explanation of solubility
H solution
-H lattice
H hydration = +
For Group II sulphates, the cations are much smaller
than the anions. The changing in size of cations does
not cause a significant change in H lattice (proportional
to 1/(r+ + r-).
However, the changing in size of cations does cause
H hydration (proportional to 1/r+ and 1/r-) to become less
exothermic, and the solubility decreases when
descending the Group.

50. SO42-
MgSO4
SrSO4

51. Explanation of solubility
H solution
-H lattice
H hydration = +
For the smaller size anions, OH-.
Down the Group, less enthalpy is required to
break the lattice as the size of cation increases.
However the change in H solution is comparatively
smaller due to the large value of 1/r- .
As a result, H solution becomes more exothermic
and the solubility increases down the Group.

52. Mg(OH)2
Sr(OH)2

53. Uses of s-block compounds
Sodium carbonate
Manufacture of glass
Water softening
Paper industry
Sodium hydrocarbonate
Baking powder
Soft drink

54. Uses of s-block compounds
Sodium hydroxide
Manufacture of soaps, dyes, paper and drugs
To make rayon and important chemicals
Magnesium hydroxide
Milk of magnesia, an antacid
Calcium hydroxide
To neutralize acids in waste water treatment
Strontium compound
Fireworks, persistent intense red flame

55. Thank You